Thermochemistry Experiment:

There is no meter or device which can measure heat directly, so calorimetry is necessarily an indirect technique (a thermometer measures temperature; not heat). In an isolated physical system under conditions of constant pressure, conservation of total energy and the First Law of Thermodynamics dictate that there can be no change in the total heat content of the system. This means that any heat generated or released by one part of the system is absorbed totally by the other parts of the system. Chemical or physical processes taking place within the system may absorb or release heat, but this must result in the entire system changing its temperature. This is the physical basis of calorimetry. Precise measurements of changes in temperature are converted into quantities of heat.

When a reaction occurs that is exothermic (gives off heat) the surrounds will appear to get warmer since the heat lost by the reaction flows into and is measured in the surroundings. Likewise when a reaction is endothermic we observe a cooling in the surroundings since the reaction absorbs heat.

One can think of heat like money. If you give money to a friend, your income is negative and that of your friend’s is positive. It’s the same with heat. When two systems are in thermal contact, if one system gains heat (positive), then the other lose some (negative). Thus,

- q reaction = + q surroundings.Where q represents the heat transferred.

Note that mathematically +q reaction = -q surroundings is identical. In this equation the negative sign simply means that what one system loses the other gains; it does not tell you if the reaction is gaining or losing heat. You must actually measure q for the reaction to determine this.

The change in temperature of an object depends on both the quantity of heat absorbed and its specific heat according to the relationship:

q = m S T

Where m is the mass of the object, S is the specific heat and T the change in temperature.

When a reaction occurs in a calorimeter the heat lost by the reaction (system) is gained by the calorimeter (surroundings) and so we get:

- q reaction = +q calorimeter = + m S T

Where m, S, and T refer to the mass, specific heat, and change in temperature of the calorimeter. Note that the q reaction is proceeded by a minus sign since the heat lost or gained by the system must have the opposite sign as the heat measured in the surroundings (remember the money analogy!).

1a. The Specific Heat of a Metal

In the first part of the experiment we will determine the specific heat of a metal. This will be accomplished by heating the metal to 100.0ºC and then adding the hot metal to our calorimeter. –q for the metal will be equal to +q for the water, thus using the temperature change of the metal and the water and their respective masses, we can find the specific heat of the metal:

-q metal = +q water

-m S T (metal) = + m S T (water)

1b. The Law of Dulong and Petit

Prior to our modern understanding of thermochemistry, Dulong and Petit discovered that about 25 joules of energy are required to raise the temperature of one mole of most metals one degree Celsius. This method can be used to approximate the molar mass of a metal:

You will use the Law of Dulong and Petit to estimate the molar mass for your metal.

2. Heats of Solution

In the second part of the experiment we will determine the heat of solution (the heat released or absorbed when a solid is dissolved) of an unknown salt.

Since,

- q reaction (dissolution) = + m S T (water)

In this experiment we will have to assume that the specific heat of the water in the calorimeter does not change significantly when the salt is added. For dilute solutions, this is an excellent assumption.

Experimental Procedure:

Determining the Specific Heat of a Metal

Fill a 400 mL beaker two-thirds full of water and begin heating it to a boil.

Obtain a calorimeter (this consist of two nested coffee cups and a lid). Insert your thermometer through the lid.

Weigh the calorimeter (including the lid and thermometer). Record the weight. Now add about 40 mL of water and then weigh the calorimeter with the water to determine the exact mass of the water added.

Weigh your entire unknown metal sample (be sure the metal is dry). Fill a large, clean, dry, loosely stoppered test-tube with your unknown metal (be careful not to break the test-tube when adding the metal sample). Wait until the water in the beaker comes to a full boil. Put the stoppered test-tube containing the metal into the beaker of boiling water. The water level in the beaker should be high enough that it covers all the metal in the tube, but not so high that water gets into the tube. You may want to use your wire test-tube holder to help support the test-tube in the beaker while heating. Heat the metal for at least 10 minutes to ensure that the metal sample reaches the same temperature as the boiling water.

Measure the temperature of the boiling water to the nearest 0.1ºC. It should be at (or close to) 100ºC. Slight variations may be due to atmospheric pressure – but if there is a significant deviation from 100ºC you may want to use a different thermometer for the remainder of the experiment.

Now begin measuring the temperature of the water in the calorimeter to the nearest 0.1ºC. Record the temperature of the water every 30 seconds for 5-7 minutes, or until the temperature remains fairly constant for a period of at least 2 minutes AND the metal has been heating for at least 10 minutes.

The following step must be performed in just a few seconds or significant heat loss to the atmosphere may occur:

One person should remove the lid and thermometer from the calorimeter and then a second person should remove the large test-tube containing the metal from the boiling water, quickly remove the stopper and transfer all of the metal to the open calorimeter. The lid and thermometer should then be replaced as quickly as possible.

Gently swirl the water and metal in the calorimeter (being careful not to break the thermometer or spill any water) and then measure the temperature of the water/metal mixiture. Make another measurement every 30 seconds for 2 minutes swirling gently between trials. Record the maximum temperature reached in the data table.

Repeat the experiment two more times, the next time using 45 mL and then 50 mL of water inside the calorimeter. Be sure to dry the metal before heating again.

Heats of Solution

Place about 50 mL of water into the calorimeter and determine the exact mass of the water as in part one.

Measure the mass of a small dry empty beaker. Measure out about 5 grams of your unknown salt into a small dry empty beaker. Record the exact mass of the salt sample and the beaker to 3 decimal places (do not subtract the mass of the beaker from this measurement).

Measure the temperature of the water in the calorimeter to the nearest 0.1ºC. Record the temperature of the water every 30 seconds for 5-7 minutes, or until the temperature remains fairly constant for a period of at least 2 minutes.

Quickly removed the lid and pour the unknown salt sample into the calorimeter. Don't worry if some of the salt remains in the beaker as this will be accounted for later. Close the calorimeter and swirl gently. Measure the temperature of the solution. Make another measurement every 30 seconds for 2 minutes swirling gently between trials. Record the maximum or minimum temperature reached in the data table.

After the reaction is complete re-measure the mass of the beaker with any remaining salt. You can subtract this from the mass of the beaker + salt to determine the mass of salt added by difference.

Repeat the experiment three times.

THERMOCHEMISTRY REPORT:

Team Name: ______

Team Members: ______

Determining the Specific Heat of a Metal:

Unknown Metal #: ______

Physical description of metal (color, shape, etc.):

Trial 1 / Trial 2 / Trial 3
Mass of empty calorimeter / N/A / N/A
Mass of calorimeter and water
Mass of water in calorimeter (calculated)
Mass of unknown metal sample
Temperature of boiling water
Average temperature of water in calorimeter after 5-7 minutes
Maximum recorded temperature of water in calorimeter after adding hot metal
∆T water (T final – T initial; calculated)
∆T metal (T final – T initial; calculated)
q water (calculated)
Specific heat of the metal in J/gºC
Estimated molar mass of the metal using
the Law of Dulong and Petit:
Average Specific Heat:
Average Estimated Molar Mass:
  1. Your metal is one of the following: Lead, Copper, Aluminum, Cobalt, Zinc, or Nickel. Look up the specific heats of these metals and then using this and your physical description of the metal identify your unknown:

Metal / Specific Heat* / Metal / Specific Heat*
Lead / Cobalt
Copper / Zinc
Aluminum / Nickel

(*) Cite Reference used: ______

Identity of your unknown: ______

Supporting data:

  1. Based on your identity of the metal above, calculate the percent errors for both your average specific heat and average molar mass:
  1. Percent error in Specific Heat:
  1. Percent error in Molar Mass:
  1. Explain the percent error you obtained for the specific heat of the metal. What experimental factors may have resulted in your error. List at least 3 factors that can explain the error you observed.

a.

b.

c.

  1. How might you improve the experiment to reduce each of the three errors you listed?

a.

b.

c.

  1. How well did the law of Dulong and Petit work for your metal? Is this a good method for determining molecular weights? Why or why not?

Heats of Solution

Unknown Salt #: ______

Trial 1 / Trial 2 / Trial 3
Mass of empty calorimeter (from part one) / N/A / N/A
Mass of calorimeter and water
Mass of water in calorimeter (calculated)
Mass of empty beaker
Mass of salt sample and beaker
Mass of beaker after pouring salt into calorimeter
Mass of salt added (calculated by difference)
Average initial temperature of water in calorimeter after 5-7 minutes.
Minimum or Maximum final temperature reached by the salt/water mixture.
∆T water (T final – T initial; calculated)
q water (joules; calculated)
∆H reaction (joules; calculated)
∆H reaction (joules/g; calculated)
∆H reaction average (joules/g)
  1. Based on your data is the reaction exothermic or endothermic (circle one)
  1. Assume you wanted to construct a commercial hot or cold pack such as the ones used for sports injuries using this salt. If the pack contains 100.-mL of water, how many grams of salt (in a separate breakable pouch) would be needed to have the temperature of the water heat to 40ºC, or cool to 5ºC (depending on if your salt solution is exothermic or endothermic).